MCAT General Chemistry Review Summary

This MCAT General Chemistry Review Summary Page is by no means an exhaustive review of MCAT General Chemistry. Our summary is only meant to highlight key points that are most helpful for the MCAT. For a list of topics for the MCAT: MCAT General Chemistry Topics List. You can also access practice questions in our Free MCAT practice test , or our many full-length MCAT practice tests.

For MCAT General Chemistry, it helps to understand vocabulary, definitions and equations. In some respects, this General Chemistry cheat sheet will minimize your need to memorize information and maximize your General Chemistry review.

Note: Unless mentioned otherwise, the following images are excerpts from the Gold Standard MCAT General Chemistry ebook.

• Mole - Atomic and Molecular Weights

  

• Categories of Chemical Reactions

  

  Note: Any reaction that does not involve the transfer of electrons (= change in oxidation numbers) qualifies as a non-redox reaction. Combination reactions qualify as non-redox reactions when all reactants and products are compounds and the oxidation numbers do not change. Decomposition reactions qualify as non-redox reactions when all reactants and products are compounds and the oxidation numbers do not change. Additional information: UNC-Chapel Hill Chemistry Fundamentals Program

  
  
  

• Oxidation Numbers, Redox Reactions, Oxidizing vs. Reducing Agents

  • Here are the general rules:
    • In elementary substances, the oxidation number of an uncombined element is zero
    • In monatomic ions the oxidation number of the elements that make up this ion is equal to the charge of the ion
    • In a neutral molecule the sum of the oxidation numbers of all the elements that make up the molecule is zero
  • Some useful oxidation numbers to memorize
    • For H: +1, except in metal hydrides where it is equal to -1
    • For O: -2 in most compounds; In peroxides (e.g. in H₂O₂) the oxidation number for O is -1, it is +2 in OF₂ and -1/2 in superoxides
    • For alkali metals: +1
    • For alkaline earth metals: +2
    • Aluminium always has an oxidation number of +3 in all its compounds
    

• Mixtures

  

• Conventional Notation for Electronic Structure

The order for filling atomic orbitals: Follow the direction of successive arrows moving from top to bottom.

• Metals, Nonmetals and Metalloids

*General Characteristics of metals, nonmetals and metalloids

Metals	Nonmetals	Metalloids
Hard and shiny	Gases or dull, brittle solids	Appearance will vary
3 or less valence electrons	5 or more valence electrons	3 to 7 valence electrons
Form + ions by losing e－	Form － ions by gaining e－	Form + and/or －ions
Good conductors of heat and electricity	Poor conductors of heat and electricity	Conduct better than nonmetals, but not as well as metals

*These are general characteristics. There are exceptions beyond the scope of the exam.

• Partial Ionic Character

  • This polar bond will also have a dipole moment given by:

    
    where q is the charge and d is the distance between these two atoms.
• Lewis Acids and Lewis Bases

  • The Lewis acid BF₃ and the Lewis base NH₃. Notice that the green arrows follow the flow of electron pairs.

    
• Valence Shell Electronic Pair Repulsions (VSEPR Models)

  • Geometry of simple molecules in which the central atom A has one or more lone pairs of electrons (= e^-)

    

    Note: dotted lines only represent the overall molecular shape and not molecular bonds. In brackets under "Molecular Geometry" is the hybridization as discussed in ORG 1.2.

    
    

    Molecular arrangement of electron pairs around a central atom A. Dotted lines only represent the overall molecular shape and not molecular bonds.

• Standard Temperature and Pressure, Standard Molar Volume
  • 0 °C (273.15 K) and 1.00 atm (101.33 kPa = 760 mmHg = 760 torr); these conditions are known as the standard temperature and pressure (STP). {Note: the SI unit of pressure is the pascal (Pa).}
  • The volume occupied by one mole of any gas at STP is referred to as the standard molar volume and is equal to 22.4 L.
• Kinetic Molecular Theory of Gases (A Model for Gases)
  • The average kinetic energy of the particles (KE = 1/2 mv²) increases in direct proportion to the temperature of the gas (KE = 3/2 kT) when the temperature is measured on an absolute scale (i.e. the Kelvin scale) and k is a constant (the Boltzmann constant).
    
    

    The Maxwell Distribution Plot

    Graham’s Law (Diffusion and Effusion of Gases)	Combined Gas Law
    Charles’ Law	Ideal Gas Law since m/V is the density (d) of the gas:
    Boyle’s Law	Partial Pressure and Dalton’s Law Of course, the sum of all mole fractions in a mixture must equal one: ΣX1 = 1
    Avogadro’s Law	The partial pressure (Pi) of a component of a gas mixture is equal to:

• Liquid Phase (Intra- and Intermolecular Forces)

  

  Van Der Waal's forces (weak) and hydrogen bonding (strong). London forces between Cl₂ molecules, dipole-dipole forces between HCl molecules and H-bonding between H₂O molecules. Note that a partial negative charge on an atom is indicated by ẟ- (delta negative), while a partial positive charge is indicated by ẟ+ (delta positive). Notice that one H₂O molecule can potentially form 4 H-bonds with surrounding molecules which is highly efficient. The preceding is one key reason that the boiling point of water is higher than that of ammonia, hydrogen fluoride, or methanol.

• Surface Tension
  • PE is directly proportional to the surface area (A)
  • PE = gA; g = surface tension
  • g = F/l; F = force of contraction of surface; l = length along surface
  
  (a) cohesive > adhesive
  
  (b) adhesive > cohesive

• Phase Changes

• Phase Diagrams

• Ions in Solution
  • Ions that are positively charged = cations; ions that are negatively charged = anions
  • Mnemonic: anions are negative ions
  • The word “aqueous” simply means containing or dissolved in water
  

  Common Anions and Cations

• Solubility Rules
  1. All salts of alkali metals are soluble.
  2. All salts of the ammonium ion are soluble.
  3. All chlorides, bromides and iodides are water soluble, with the exception of Ag⁺, Pb²⁺, and Hg₂²⁺.
  4. All salts of the sulfate ion (SO₄^2-) are water soluble with the exception of Ca²⁺, Sr²⁺, Ba²⁺, and Pb²⁺.
  5. All metal oxides are insoluble with the exception of the alkali metals and CaO, SrO and BaO.
  6. All hydroxides are insoluble with the exception of the alkali metals and Ca²⁺, Sr²⁺, Ba²⁺.
  7. All carbonates (CO₃^2-), phosphates (PO₄^3-), sulfides(S^2-) and sulfites (SO₃^2-) are insoluble, with the exception of the alkali metals and ammonium.

• The First Law of Thermodynamics

  
  • heat absorbed by the system: Q > 0
  • heat released by the system: Q < 0
  • work done by the system on its surroundings: W > 0
  • work done by the surroundings on the system: W < 0

• Temperature Scales

  

• State Functions

  • W can be determined experimentally by calculating the area under a pressure-volume curve

    	Work	Heat	Changes in internal energy
    1st tranf.	w	0	-w
    2nd transf.	W = w +q	q	-w

• Heat of Reaction: Basic Principles
  • A reaction during which heat is released is said to be exothermic (ΔH is negative).

  • If a reaction requires the supply of a certain amount of heat it is endothermic (ΔH is positive).

    

• Bond Dissociation Energies and Heats of Formation

  

• Calorimetry

  

• The Second Law of Thermodynamics

  • For any spontaneous process, the entropy of the universe increases which results in a greater dispersal or randomization of the energy (ΔS > 0).

• Entropy

  

• Free Energy

  
  • A reaction carried out at constant pressure is spontaneous if: ΔG < 0
  • It is not spontaneous if: ΔG > 0
  • It is in a state of equilibrium (reaction spontaneous in both directions) if: ΔG = 0

• Dependence of Reaction Rates on Concentration of Reactants

  
  • [ ] is the concentration of the corresponding reactant in moles per liter
  • k is referred to as the rate constant
  • m is the order of the reaction with respect to A
  • n is the order of the reaction with respect to B
  • m+n is the overall reaction order

• Dependence of Reaction Rates upon Temperature

  

  Potential Energy Diagrams: Exothermic vs. Endothermic Reactions

• Catalysis